Carbon - Real-life applications

O
RGANIC
C
HEMISTRY

We have stated that carbon forms tetravalent bonds, and makes multiple
bonds with a single atom. In addition, we have mentioned the fact that
carbon forms long chains of atoms and varieties of shapes. But how
does it do these things, and why? These are good questions, but not
ones we will attempt to answer here. In fact, an entire branch of
chemistry is devoted to answering these theoretical questions, as well
as to determining solutions to a host of other, more practical
problems.

Organic chemistry is the study of carbon, its compounds, and their
properties. (There are carbon-containing compounds that are not
considered organic, however. Among these are oxides such as carbon
dioxide and monoxide; as well as carbonates, most notably calcium
carbonate.) At one time, chemists thought that "organic"
was synonymous with "living," and even as recently as
the early nineteenth century, they believed that organic substances
contained a supernatural "life force." Then, in 1828,
German chemist Friedrich Wöhler (1800-1882) cracked the code that
distinguished the living from the nonliving, and the organic from the
inorganic.

Wöhler took a sample of ammonium cyanate (NH
4
OCN), and by heating it, converted it into urea (H
2
N-CO-NH
2
), a waste product in the urine of mammals. In other words, he had
turned an inorganic material into a organic one, and he did so, as he
observed, "without benefit of a kidney, a bladder, or a
dog." It was almost as though he had created life. In fact,
what Wöhler had glimpsed—and what other scientists who
followed came to understand, was this: what separates the organic from
the inorganic is the manner in which the carbon chains are arranged.

Ammonium cyanate and urea have exactly the same numbers and
proportions of atoms, yet they are different compounds. They are thus
isomers: substances which have the same formula, but are different
chemically. In urea, the carbon forms an organic chain, and in
ammonium cyanate, it does not. Thus, to reduce the specifics of
organic chemistry even further, it can be said that this area of the
field constitutes the study of carbon chains, and ways to rearrange
them in order to create new substances.

Rubber, vitamins, cloth, and paper are all organically based compounds
we encounter in our daily lives. In each case, the material comes from
something that once was living, but what truly makes these substance
organic in nature is the common denominator of carbon, as well as the
specific arrangements of the atoms. We have organic chemistry to thank
for any number of things: aspirins and all manner of other drugs;
preservatives that keep food from spoiling; perfumes and toiletries;
dyes and flavorings, and so on.

A
LLOTROPES OF
C
ARBON

GRAPHITE.

Carbon has several allotropes—different versions of the same
element, distinguished by molecular structure. The first of these is
graphite, a soft material with an unusual crystalline structure.
Graphite is essentially a series of one-atom-thick sheets of carbon,
bonded together in a hexagonal pattern, but with only very weak
attractions between adjacent sheets. A piece of graphite is thus like
a big, thick stack of carbon paper: on the one hand, the stack is
heavy, but the sheets are likely to slide against one another.

Actually, people born after about 1980 may have little experience with
carbon paper, which was gradually phased out as photocopiers became
cheaper and more readily available. Today, carbon paper is most often
encountered when signing a credit-card receipt: the signature goes
through the graphite-based backing of the receipt, onto a customer
copy.

In such a situation, one might notice that the copied image of the
signature looks as though it were signed in pencil. This is not
surprising, considering that pencil "lead" is, in fact,
a mixture of graphite, clay, and wax. In ancient times, people did
indeed use lead—the heaviest member of Group 4, the
"carbon family"—for writing, because it left gray
marks on a surface. Lead, of course, is poisonous, and is not used
today in pencils or in most applications that would involve prolonged
exposure of humans to the element. Nonetheless, people still use the
word "lead" in reference to pencils, much as they still
refer to a galvanized steel roof with a zinc coating as a "tin
roof."

In graphite the atoms of each "sheet" are tightly bonded
in a hexagonal, or six-sided, pattern, but the attractions between the
sheets are not very strong. This makes it highly useful as a lubricant
for locks, where oil would tend to be messy. A good conductor of
electricity, graphite is also utilized for making high-temperature
electrolysis cells. In addition, the fact that graphite resists
temperatures of up to about 6,332°F (3,500°C) makes it
useful in electric motors and generators.

DIAMOND.

The second allotrope of carbon is also crystalline in structure. This
is diamond, most familiar in the form of jewelry, but in fact widely
applied for a number of other purposes. According to the Moh scale,
which measures the hardness of minerals, diamond is a 10—in
other words, the hardest type of material. It is used for making
drills that bore through solid rock; likewise, small diamonds are used
in dentists' drills for boring through the ultra-hard enamel on
teeth.

Neither diamonds nor graphite are, in the strictest sense of the term,
formed of molecules. Their arrangement is definite, as with a
molecule, but their size is not: they simply form repeating patterns
that seem to stretch on forever. Whereas graphite is in the form of
sheets, a diamond is basically a huge "molecule"
composed of carbon atoms strung together by covalent bonds. The size
of this "molecule" corresponds to the size of the
diamond: a diamond of 1 carat, for instance, contains about 10
22
(10,000,000,000,000,000,000,000 or 10 billion billion) carbon atoms.

The diamonds used in industry look quite different from the ones that
appear in jewelry. Industrial diamonds are small, dark, and cloudy in
appearance, and though they have the same chemical properties as
gem-quality diamonds, they are cut with functionality (rather than
beauty) in mind. A diamond is hard, but brittle: in other words, it
can be broken, but it is very difficult to scratch or cut a
diamond—except with another diamond.

The cutting of fine diamonds for jewelry is an art, exemplified in the
alluring qualities of such famous gems as the jewels in the British
Crown or the infamous Hope Diamond in Washington, D.C.'s
Smithsonian Institution. Such diamonds—as well as the diamonds
on an engagement ring—are cut to refract or bend light rays,
and to disperse the colors of visible light.

BUCKMINSTERFULLERENE.

Until 1985, carbon was believed to exist in only two crystalline
forms, graphite and diamond. In that year, however, chemists at Rice
University in Houston, Texas, and at the University of Sussex in
England, discovered a third variety of carbon—and later jointly
received a Nobel Prize for their work. This "new" carbon
molecule composed of 60 bonded atoms in the shape of what is called a
"hollow truncated icosahedron." In plain language, this
is rather like a soccer ball, with interlocking pentagons and
hexagons. However, because the surface of each geometric shape is
flat, the "ball" itself is not a perfect sphere. Rather,
it describes the shape of a geodesic dome, a design created by
American engineer and philosopher R. Buckminster Fuller (1895-1983).

There are other varieties of buckminsterfullerene molecules, known as
fullerenes. However, the 60-atom shape, designated as
60
C, is the most common of all fullerenes, the result of condensing
carbon slowly at high temperatures. Fullerenes potentially have a
number of applications, particularly because they exhibit a whole
range of electrical properties: some are insulators, while some are
conductors, semiconductors, and even superconductors. Due to the high
cost of producing fullerenes artificially, however, the ways in which
they are applied remain rather limited.

AMORPHOUS CARBON.

There is a fourth way in which carbon appears, distinguished from the
other three in that it is amorphous, as opposed to crystalline, in
structure. An example of amorphous carbon is carbon black, obtained
from smoky flames and used in ink, or for blacking rubber tires.

Though it retains some of the microscopic structures of the plant
cells in the wood from which it is made, charcoal—wood or other
plant material that has been heated without enough air present to make
it burn—is mostly amorphous carbon. One form of charcoal is
activated charcoal, in which steam is used to remove the sticky
products of wood decomposition. What remains are porous grains of pure
carbon with enormous microscopic surface areas. These are used in
water purifiers and gas masks.

A
GEODESIC DOME IN
M
ONTREAL
, Q
UEBEC
, C
ANADA
. S
UCH DOMES ARE NAMED AFTER THEIR DESIGNER
, R. B
UCK
-
MINSTER
F
ULLER
,
AND EVENTUALLY PROVIDED THE NAME FOR THE CARBON MOLECULES KNOWN
AS BUCKMINSTER
-
FULLERENES
.

(Lee Snider/Corbis

.
Reproduced by permission.)

Coal and coke are particularly significant varieties of amorphous
carbon. Formed by the decay of fossils, coal was one of the first
"fossil fuels" (for example, petroleum) used to provide
heat and power for industrial societies. Indeed, when the words
"industrial revolution" are mentioned, many people
picture tall black smokestacks belching smoke from coal fires.
Fortunately—from an environmental standpoint—coal is not
nearly so widely used today, and when it is (as for instance in
electric power plants), the methods for burning it are much more
efficient than those applied in the nineteenth century.

Actually, much of what those smokestacks of yesteryear burned was
coke, a refined version of coal that contains almost pure carbon.
Produced by heating soft coal in the absence of air, coke has a much
greater heat value than coal, and is still widely used as a reducing
agent in the production of steel and other alloys.

C
ARBON
D
IOXIDE

Carbon forms many millions of compounds, some families of which will
be discussed below. Two others, formed by the bonding of carbon atoms
with oxygen atoms, are of particular significance. In carbon dioxide,
a single carbon joins with two oxygens to produce a gas essential to
plant life. In carbon monoxide (CO), a single oxygen joins the carbon,
creating a toxic—but nonetheless important—compound.

The first gas to be distinguished from ordinary air, carbon dioxide is
an essential component in the natural balance between plant and animal
life. Animals, including humans, produce carbon dioxide by breathing,
and humans further produce it by burning wood and other fuels. Plants
use carbon dioxide when they store energy in the form of food, and
they release oxygen to be used by animals.

DISCOVERY.

Flemish chemist and physicist Johannes van Helmont (1579-1644)
discovered in 1630 that air was not, as had been thought up to that
time, a single element: it contained a second substance, produced in
the burning of wood, which he called "gas sylvestre."
Thus he is recognized as the first scientist to note the existence of
carbon dioxide.

More than a century later, in 1756, Scottish chemist Joseph Black
(1728-1799) showed that carbon dioxide—which he called
"fixed air"—combines with other chemicals to form
compounds. This and other determinations Black made concerning carbon
dioxide led to enormous

S
OFT DRINKS ARE MADE POSSIBLE BY THE USE OF CARBONATED WATER
.

(Sergio Dorantes/Corbis

.
Reproduced by permission.)

progress in the discovery of gases by various chemists of the late
eighteenth century.

By that time, chemists had begun to arrive at a greater degree of
understanding with regard to the relationship between plant life and
carbon dioxide. Up until that time, it had been believed that plants
purify the air by day, and poison it at night. Carbon dioxide and its
role in the connection between animal and plant life provided a much
more sophisticated explanation as to the ways plants
"breathe."

USES.

Around the same time that Black made his observations on carbon
dioxide, English chemist Joseph Priestley (1733-1804) became the first
scientist to put the chemical to use. Dissolving it in water, he
created carbonated water, which today is used in making soft drinks.
Not only does the gas add bubbles to drinks, it also acts as a
preservative.

Though the natural uses of carbon dioxide are by far the most
important, the compound has numerous industrial and commercial
applications. Used in fire extinguishers, carbon dioxide is ideal for
controlling electrical and oil fires, which cannot be put out with
water. Heavier than air, carbon dioxide blankets the flames and
smothers them.

In the solid form of dry ice, carbon dioxide is used for chilling
perishable food during transport. It is also one of the only compounds
that experiences sublimation, or the instantaneous transformation of a
solid to a gas without passing through an intermediate liquid state,
at conditions of ordinary pressure and temperature. Dry ice has often
been used in movies to generate "mists" or
"smoke" in a particular scene.

C
ARBON
M
ONOXIDE

During the late eighteenth century, Priestley discovered a
carbon-oxygen compound different from carbon dioxide: carbon monoxide.
Scientists had actually known of this toxic gas, released in the
incomplete combustion of wood, from the Middle Ages onward, but
Priestley was the first to identify it scientifically.

Industry uses carbon monoxide in a number of ways. By blowing air
across very hot coke, the result is producer gas, which, along with
water gas (made by passing hot steam over coal) is an important fuel.
Producer gas constitutes a 6:1:18 mixture of carbon monoxide, carbon
dioxide, and nitrogen, while water gas is 40% carbon monoxide, 50%
hydrogen, and 10% carbon dioxide and other gases.

Not only are producer and water gas used for fuel, they are also
applied as reducing agents. Thus, when carbon monoxide is passed over
hot iron oxides, the oxides are reduced to metallic iron, while the
carbon monoxide is oxidized to form carbon dioxide. Carbon monoxide is
also used in reactions with metals such as nickel, iron, and cobalt to
form some types of carbonyls.

Carbon monoxide—produced by burning petroleum in automobiles,
as well as by the combustion of wood, coal, and other
carbon-containing fuels—is extremely hazardous to human health.
It bonds with iron in hemoglobin, the substance in red blood cells
that transports oxygen throughout the body, and in effect fools the
body into thinking that it is receiving oxygenated hemoglobin, or
oxyhemoglobin. Upon reaching the cells, carbon monoxide has much less
tendency than oxygen to break down, and therefore it continues to
circulate throughout the body. Low concentrations can cause nausea,
vomiting, and other effects, while prolonged exposure to high
concentrations can result in death.

C
ARBON AND THE
E
NVIRONMENT

Carbon is released into the atmosphere by one of three means: cellular
respiration; the burning of fossil fuels; and the eruption of
volcanoes. When plants take in carbon dioxide from the atmosphere,
they combine this with water and manufacture organic compounds using
energy they have trapped from sunlight by means of
photosynthesis—the conversion of light to chemical energy
through biological means. As a by-product of photosynthesis, plants
release oxygen into the atmosphere.

In the process of undergoing photosynthesis, plants produce
carbohydrates, which are various compounds of carbon, hydrogen, and
oxygen essential to life. The other two fundamental components of a
diet are fats and proteins, both carbon-based as well. Animals eat the
plants, or eat other animals that eat the plants, and thus incorporate
the fats, proteins, and sugars (a form of carbohydrate) from the
plants into their bodies. Cellular respiration is the process whereby
these nutrients are broken down to create carbon dioxide.

Photosynthesis and cellular respiration are thus linked in what is
known as the carbon cycle. Cellular respiration also releases carbon
into the atmosphere through the action of decomposers—bacteria
and fungi that feed on the remains of plants and animals. The
decomposers extract the energy in the chemical bonds of the
decomposing matter, thus releasing more carbon dioxide into the
atmosphere.

When creatures die and are buried in such a way that they cannot be
reached by decomposers—for instance, at the bottom of the
ocean, or beneath layers of rock—the carbon in their bodies is
eventually converted to fossil fuels, including petroleum, natural
gas, and coal. The burning of fossil fuels releases carbon (both
monoxide and dioxide) into the atmosphere.

Because the rate of such burning has increased dramatically since the
late nineteenth century, this has raised fears that carbon dioxide in
the atmosphere may create a greenhouse effect, leading to global
warming. On the other hand, volcanoes release tons of carbon into the
atmosphere regardless of whether humans burn fossil fuels or not.

R
ADIOCARBON
D
ATING

Radiocarbon dating is used to date the age of charcoal, wood, and
other biological materials. When an organism is alive, it incorporates
a certain ratio of carbon-12 in proportion to the amount of the
radioisotope (that is, radioactive isotope) carbon-14 that it receives
from the atmosphere. As soon as the organism dies, however, it stops
incorporating new carbon, and the ratio between carbon-12 and
carbon-14 will begin to change as the carbon-14 decays to form
nitrogen-14.

Carbon-14 has a half-life of 5,730 years, meaning that it takes that
long for half the isotopes in a sample to decay to nitrogen-14.
Therefore a scientist can use the ratios of carbon-12, carbon-14, and
nitrogen-14 to guess the age of an organic sample. The problem with
radiocarbon dating, however, is that there is a good likelihood the
sample can become contaminated by additional carbon from the soil.
Furthermore, it cannot be said with certainty that the ratio of
carbon-12 to carbon-14 in the atmosphere has been constant throughout
time.